Short Answer

Molecular orbital theory (MOT) explains how atoms combine to form molecules by allowing their atomic orbitals to mix and form new orbitals called molecular orbitals. These orbitals belong to the whole molecule, not to individual atoms. Electrons fill these molecular orbitals according to the same rules used for atomic orbitals.

This theory helps describe important properties like bond order, magnetic nature, stability, and energy levels of molecules. MOT gives a more accurate picture of bonding than simple valence bond theory, especially for molecules like O₂, which shows paramagnetism.

Detailed Explanation :

Molecular Orbital Theory

Molecular orbital theory (MOT) is a modern and advanced way of describing chemical bonding in molecules. Instead of assuming that electrons are shared between two atoms as in valence bond theory, MOT states that when atoms come close together, their atomic orbitals combine to form new orbitals that spread over the entire molecule. These new orbitals are called molecular orbitals.

Molecular orbitals can hold electrons from all the atoms in the molecule. They determine the energy, magnetic behavior, bond strength, and stability of the molecule. MOT explains many observations that older theories cannot, such as the magnetic nature of oxygen and bond order in molecules like N₂, O₂, and F₂.

1. Formation of Molecular Orbitals

When two atomic orbitals combine, they form two new molecular orbitals:

1. Bonding molecular orbital (lower energy)
2. Antibonding molecular orbital (higher energy)

The combination depends on the phase and symmetry of the atomic orbitals.

• When orbitals combine constructively (same phase), a bonding orbital forms.
• When orbitals combine destructively (opposite phases), an antibonding orbital forms.

Bonding orbitals increase electron density between nuclei, stabilizing the molecule.
Antibonding orbitals decrease electron density and destabilize the molecule.

2. Types of Molecular Orbitals

(a) Sigma (σ) Molecular Orbitals

Form when orbitals overlap along the internuclear axis.
Examples:

• s–s overlap
• s–p overlap
• p–p overlap (end-to-end)

Sigma bonding orbitals are stronger and more stable.

(b) Pi (π) Molecular Orbitals

Form when p orbitals overlap sideways.
Pi bonds are weaker and occur in molecules with double or triple bonds.

3. Filling of Molecular Orbitals

Electrons fill molecular orbitals using principles similar to atomic orbitals:

• Aufbau principle: fill lower energy orbitals first
• Pauli exclusion principle: each orbital holds two electrons with opposite spins
• Hund’s rule: electrons fill orbitals singly before pairing

This distribution determines the molecule’s properties, such as magnetic behavior.

4. Bond Order According to MOT

Bond order tells whether a molecule is stable and how strong the bond is.

Examples:

• N₂ has a bond order of 3 → very strong triple bond
• O₂ has a bond order of 2 → double bond
• He₂ has a bond order of 0 → does not exist as a stable molecule

Bond order increases as bonding electrons increase or antibonding electrons decrease.

5. Magnetic Properties Explained by MOT

MOT correctly explains magnetic behavior:

• Molecules with unpaired electrons in molecular orbitals are paramagnetic.
• Molecules with no unpaired electrons are diamagnetic.

Example:
O₂ is paramagnetic because it has two unpaired electrons in π* molecular orbitals.
Simple theories cannot explain this, but MOT does.

6. Energy Level Diagrams of Diatomic Molecules

Different molecules follow different orbital-filling patterns:

• For molecules like B₂, C₂, and N₂, the π orbitals have lower energy than σ2p.
• For O₂ and F₂, the σ2p orbital is lower than the π orbitals.

This order explains differences in bond order and magnetic behavior among molecules.

7. Advantages of Molecular Orbital Theory

MOT explains several properties that valence bond theory cannot:

• Magnetic properties (paramagnetism of O₂)
• Variations in bond order and bond energy
• Delocalization of electrons over the whole molecule
• Stability of ions like O₂⁺ and O₂⁻
• Existence of molecules not predicted by VB theory (like He₂⁺)

MOT gives a clearer, more mathematical description of bonding.

8. Electron Delocalization

One of the major strengths of MOT is its ability to describe delocalized electrons.
Electrons in molecular orbitals are not confined to one bond or one atom; they are spread across the entire molecule.

This helps explain:

• Aromaticity
• Resonance behavior
• Stability of conjugated systems

For example, benzene’s stability arises from delocalized π electrons, described well by MOT.

Conclusion

Molecular orbital theory states that atomic orbitals combine to form molecular orbitals that belong to the whole molecule. Electrons fill these orbitals according to standard rules, and their arrangement determines bond order, stability, and magnetic behavior. MOT successfully explains properties that older theories cannot, making it a powerful and modern method for understanding chemical bonding.